Chemistry Review Flashcards

This document provides a review of key concepts in chemistry, including the periodic table, atomic structure, molecular representation, and various types of chemical bonds. It is designed as a preparatory resource to help students better understand the material presented during the year. ## Key Concepts ### Atomic Structure * **Atomic Composition**: Atoms consist of a nucleus (positively charged) and an electron cloud (negatively charged). * Atomic Number (Z): Number of protons in the nucleus. * Mass Number (A): Total number of protons and neutrons in the nucleus. * Isotopes: Atoms with the same atomic number but different mass numbers. ### Periodic Table * Rows (Periods): Elements in the same row have the same number of electron shells. * Columns (Families): Elements in the same column have similar valence electron configurations and chemical properties. * Key groups include alkali metals, alkaline earth metals, halogens, noble gases, and transition metals. ### Molecular Representation * Different Notations: Molecules can be represented using various formulas, including brute formula, developed formula, semi-developed formula, Lewis formula and Cram representation. ### Electronic Configuration * Definition: The arrangement of electrons in the various orbitals of an atom. * Rules for Filling Orbitals: * Klechkowski's Rule: Orbitals are filled in the order of increasing (n + l) values. * Pauli Exclusion Principle: No two electrons in an atom can have the same set of quantum numbers. * Hund's Rule: Electrons fill orbitals and subshells singly before pairing up. * Exceptions: Chromium and copper have irregular electronic configurations to achieve greater stability. * Valence Electrons: Electrons in the outermost shell (or subshells) that participate in chemical bonding. ### Periodic Trends and Properties * Electronegativity: The ability of an atom to attract electrons in a chemical bond. Increases across a period and decreases down a group. ### Chemical Bonding * Covalent Bonds: Formed by sharing electrons between atoms. * Lewis Structures: Diagrams representing the valence electrons in a molecule. * Duet and Octet Rules: Atoms tend to gain, lose, or share electrons to achieve a stable electron configuration (like noble gases). * Single, double, and Triple Bonds: Bonds involving one, two, or three pairs of electrons. * VSEPR Theory: Predicts the geometry of molecules based on minimizing electron pair repulsion. * Molecular Orbitals (MOs): Formed by combining atomic orbitals. * Bonding and Antibonding Orbitals: Bonding orbitals are lower in energy and contribute to bonding, while antibonding orbitals are higher in energy and weaken the bond * Sigma (σ) and Pi (π) Bonds: Sigma bonds are formed by axial overlap, while pi bonds are formed by lateral overlap. * Hybridization: Mixing of atomic orbitals to form new hybrid orbitals with different shapes and energies. * sp, sp2, and sp3 Hybridization: Different types of hybridization result in different molecular geometries. * Conjugation and Mesomerism: Delocalization of electrons in molecules with alternating single and multiple bonds, leading to resonance structures. ### Mseomeric Effect * **Definition**: Delocalization of electrons in conjugated molecules. * Resonance Structures: Different Lewis structures that contribute to the overall electronic structure of a molecule. ### Polarization * Dipole Moment: Measure of the polarity of a molecule. * Inductive Effect: Polarization of sigma bonds due to electronegativity differences. ## Key Takeaways

Understanding these fundamental concepts is crucial for success.